chemistry ii chapter 5 thermochemistry

Chemistry IIChapter 5

Thermochemistry

1. The study of energy and its transformations is known as

_thermodynamics__.

2. __Thermochemistry_____ considers the relationships between chemical reactions and the energy changes involving heat.

 Nature of Energy

•The magnitude of the kinetic energy, Ek, of an object depends on its __mass__ (m) and __velocity___(v)

•  

Formula for kinetic energy:

•Ek = ½ mv2

4. Potential energy is the energy an object has by virtue of its ___position__ relative to other objects.

Types of potential energy

• gravitational•  

• elastic•  

• chemical•  

• electrical

5.__Electrostatic Forces_________ results from the interactions between charged particles.

The formula for electrostatic force is:

•Ee1 = kQ1Q2/ d

k is the constant of proportionality (8.99 x 109 J m /C2)

- C is the _coulomb______, the unit for electrical charge.- Q1 and Q2 are the _charges__ on the two particles involved

- d is the distance between the particles

8. When the charges are the same, the electrostatic force, Ek, is ______positive_______ and the charges __repel____.

9. When the charges are opposite the electrostic force, Ek, is ____negative_____ and the charges __attract____.• 

•The energy of a substance is due to

•1.) The thermal (kinetic) energy of the molecules.

•2.) The potential energy of the bonds.

11. The SI unit for energy is the __joule____ (J) named in honor of ___James Joule____ ( 1818 -1889) – a British Scientist who investigated heat and work.

•12. A __calorie_______ was originally defined as the amount of energy required to raise the temperature of 1gram of water from 14.5 to 15.5 degrees Celsius

13. Today energy is defined in terms of the joule:

1 calorie = 4.184 joules

14. When we use thermodynamics to analyze energy changes, we focus our attention on a limited and

well defined part of the universe.

a. The portion we single out to look at is called the __system___.b. Everything else is the _surroundings____.

c. A ___closed system_____ can exchange energy but not material with its surroundings.

Transferring Energy: Work and HeatEnergy is transferred in two ways:1.) To cause the motion of an object against a force 2.) To cause a temperature change

15. The energy used to cause an object to move against a force is called __work_____. 

•W= F x D

16. The other way that energy is transferred is __heat__- the energy transferred by a hotter object to a colder one.

17. _Energy______ is the capacity to do work or to transfer heat.

Section 2The First Law of Thermodynamics

18. The First Law of Thermodynamics – 

Energy is conserved

19. The _Internal_____ __energy___ of a system is the sum of all of the kinetic and potential energy in the components of the system.

•20. We represent the internal energy with the symbol _E__.

21. We can never really know the internal energy of a system, E, but we can find the ____Change____ in E (ΔE) •ΔE = Ef - Ei

22. A __Positive_____ value for ΔE results when Efinal> Einitial which means that energy has been put into the system ( __Endothermic___)

23. A ___negative__ value for ΔE results when Efinal< Einitial which means that energy has been lost to the surroundings ( _exothermic___)

Energy DiagramsExothermic * Endothermic

24. When a system undergoes any chemical or physical change, the change in internal energy ΔE is given by the heat, q, added to or released by the system plus the work done on or by the system, w.Formula:ΔE = q + w

When heat, q, is added to a system q is positiveWhen work, w, is done on a system it has a positive value

Sample exercise 5.3The hydrogen and oxygen gases in the cylinderillustrated in figure 5.3 are ignited. As the reaction occurs, the system loses 1150 J of heat to the surroundings. The reaction also causes the pistion to rise as the hot gases expand. The expanding gas does 480 J of work on the surroundings as it pushes against the atmosphere. What is the change in the internal energy of the system ?

Endothermic – heat flows into the system

Exothermic- heat flows out of the system

State FunctionsAlthough we cannot know exactly what the internal energy, E, of a system, it does have a definite value, Δ E can be determined.

26. Internal energy is a __state___ _function___. 

A state function is a property of a system that is determined by specifying the conditions of state.

27. The value of a state function depends only on its _present____ __conditon__ and not on the particular ___history___ of the sample.

Internal energy, E, a state function. Caption: E depends only on the present state of the system and not on the path by which

it arrived at that state. The internal energy of 50 g of water at 25 °C is the same whether the water is cooled from a higher temperature to 25 °C or warmed from a lower temperature to 25 °C.

28. B/c internal energy, E, is a state function Δ E depends only on the __initial_____ and _final___ states and not the change that occurred.

29.)Internal energy is a state function, but q+w is not a state function.

Section 3: Enthalpy

•30.)Most commonly the only kind of work produced by chemical change is _____mechanical_ work.

•31.)Usually reactions are carried out at __constant___atmospheric pressure.

For example: Consider the following equationZn + 2 HCl → ZnCl2 + H2

 Net IonicZn + 2H+ → Zn+2 + H2

32.) As the hydrogen gas is produced it does positive__ against the atmosphere.

We can better see this work if the reaction takes place in a reaction vessel like the one below.

33.) This kind of work is called _Pressure__ - __volume__ work (P-V Work).

• Formula:•W= -PΔV

•34. When the change in volume is positive the work done is ___negative__.

•35. The thermodynamic function called ___enthalpy___ accounts for the heat flow in chemical changes occurring at constant pressure when no forms of work are performed other than P-V work.

Enthalpy equals the internal energy plus the product of the pressure and the change in volume of the system.

•36.) The symbol for enthalpy is _____H_____.

Formula for enthalpy: H = E + PV

•Enthalpy (H) equals the total internal energy ( E) and pressure volume work.

37.) Enthalpy is a __state____ function because internal energy, E, pressure and volume are all state functions.

ΔH = ΔE + ΔPVW = - PΔV-W = P ΔVΔH = ΔE + ΔPV= q + w-w=q ( at constant pressure)ΔH = qp

Change in enthalpy equals the heat gained or lost at constant pressure.It is easier to measure the change in temperature than the change in internal energy.

•38.) When ΔH is ____positive____ (qp is positive) the system has gained heat from the surroundings.

39.) When ΔH is positive the reaction is _endothermic____.

40.) When ΔH is negative heat is ___lost___to the surroundings and the reaction is __exothermic___. 

41. Because H is a state function, ΔH depends only on the __initial__ and the __final____ states of the system.

• In section two we learned that q is not a state function. This is not a contradiction of ΔH = qp . Enthalpy change, ΔH, only is equal to the heat released or absorbed at __constant__ pressure.

Section 4: Enthalpies of Reaction

ΔH= ΔHfinal + ΔHinitial

ΔH= ΔHproducts + ΔHreactants

42. The enthalpy change that accompanies a reaction is called the __enthalpy of reaction_ or heat of reaction (ΔHrxn)(rxn is often used as an abbreviation for reaction)

Consider the following reaction:2H2 (g) + O2 (g) → 2H2O (g) ΔH= -483.6 KJ

• 43. When the above reaction is controlled so that exactly 2 moles of hydrogen gas react with exactly one mole of oxygen gas, 483.6 kilojoules of heat is released. Change in enthalpy (ΔH) is negative so the reaction is ___exothermic___.

44. Balanced chemical equations that show the associated enthalpy changes are called __thermochemical____ equations.

45. The enthalpy change accompanying a reaction can also be represented in an enthalpy _diagram_____.

46. In an ___exothermic_____ reaction the enthalpy of the products will be lower than the reactants.

47. In an ___endothermic____ reaction the enthalpy of the products will be higher than the

reactants

The following guidelines are helpful when using thermochemical equations and enthalpy diagrams.

• 48.)Enthalpy is an extensive property.• a.)_Extensive___ properties depend on

the amount of substance present (i.e., mass, volume, amount of heat)

• b.)_Intensive____ properties do not depend on the amount of substance present ( i.e., density, melting point, boiling point

49.) ΔH is directly dependent on the __amount___ of reactant consumed.

For example,If when 2 moles of hydrogen gas react with 1 mole of oxygen gas 493.6 KJ of heat is released.2H2(g) + O2(g) 2H2O(g) ΔH= -493.6 KJThen when 4 moles of hydrogen gas react with 2 moles of oxygen gas 985.2 KJ are released.

Sample Example 5.5 How much heat is released when 4.5 g of methane gas is burned in a constant pressure system?CH4(g) + 2O2(g) CO2(g) + 2H2O(l) ΔH= - 890 KJ

Practice ExerciseHydrogen Peroxide can decompose to water and oxygen by the following reaction:2H2O2 2H2O(l) + O2(g) Δ H = -196 KJCalculate the value of q when 5.00 grams of H2O2(l) decomposes at constant temperature.

50.) The enthalpy change for a reaction is equal in magnitude but opposite in sign to ΔH for the reverse reaction.

For example: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH= -890 KJCO2(g) + 2H2O(l) → CH4(g) + 2O2(g) ΔH= __890_ KJ

51.)When we reverse a reaction, the roles of reactants and products are reversed.

52.)The enthalpy change for a reaction depends on the state of the reactants and products.i.e., if the product of combustion of methane were gaseous H2O instead of liquid H2O, the ΔH would be= -802 KJ instead of -890 KJ.

53.)It is very important to specify the ___states____ of reactants and products in thermochemical equations.

Definitions

•54.) ___Enthalpy______ (H) A quantity defined by the relationship H = E + PV ; the enthalpy change, ΔH, for a reaction that occurs at constant pressure is the heat evolved or absorbed in the reaction; ΔH=qp

• 55)Enthalpy of formation(Hf)-• the enthalpy change that accompanies the formation of a substance from the most stable forms of its component elements.

•56.)Enthalpy of reaction (Hrxn)•The enthalpy change associated with a chemical reaction.

Using Enthalpy as Guide

• 57.) A process that is thermodynamically favored to happen is called a ___spontaneous______ process.

• 58.) By “spontaneous” we don’t mean that the reaction will form products without intervention. Most often some amount of energy must be imparted to get the process started.

•  

•59.)The enthalpy change of a reaction does give an indication as to whether the reaction is likely to be __spontaneous__.

• 60.)As a general rule, when the enthalpy change is _large_ it is the dominant factor in determining spontaneity.

•61.)When ΔH is large and negative the reaction tends to be __spontaneous_____.

•62.)Reactions for which ΔH is large and positive tend to be spontaneous in the __opposite___ direction

Section 5 : Calorimetry

•63.)The measurement of heat flow is _calorimetry_______

• 64.)An apparatus used to measure heat flow is a _________________.

Heat Capacity and Specific Heat

•65.)The ___heat_ _capacity__ of an object is the amount of heat required to raise the temperature by 1K or 1⁰C. ( one degree Kelvin is equal to one degree _celsius___) 1K = 1⁰C

•66.)The greater the heat capacity the _greater___ the heat required to produce a given rise in temperature.

•67.)The heat capacity of 1 mole of a substance is called its molar heat capacity.

• 68.)The heat capacity of 1 gram of a substance is called its specific heat capacity.

• 69.)Formula for specific heat capacity: • C =q/mΔT

• C – specific heat• q-heat• m-mass• ΔT – change in temp

•70.)The specific heat capacity for water is much __higher__ than for other substances. Water can _absorb/release___ much heat with little change in its temperature__.

•71.)The high specific heat of water is essential for the maintenance of our ____climate______

Sample exercise 5.6

• How much heat is needed to warm 250 g of water (about 1 cup) from 22ᵒC (about room temperature) to near its boiling point, 98 ⁰C? The specific heat of water is 4.18 J/g-K.

•  • What is the molar heat capacity of water?

 Practice Exercise

• Large beds of rock are used in solar heated homes to store heat. Assume that the specific heat of the rocks is 0.82J/g-K. Calculate the quantity of heat absorbed by 50 kg of rocks if their temperature increases by 12 ᵒC.

•  • What temperature change would these rocks

undergo if they emitted 450 KJ of heat?

 Constant Pressure Calorimetry

•72.)In calorimetry problems we will assume that the heat produced by the reaction, qrxn is entirely ____absorbed_____ by the solution; it does not escape the calorimeter.

• 73.)For an exothermic reaction, heat is _produced/released___by the reaction and absorbed___ by the solution.

•74.)For an endothermic reaction opposite is true- heat is __absorbed______ by the reactants and is _lost____ by the solution.

• qsoln = -q rxn

•qsoln = (specific heat of soln) x ( gram of soln)x ΔT)= -q rxn

•75.) For dilute aqueous solutions, the specific heat of the solution will be the same as water

( 4.18 J/g-k)

Sample Exercise 5.7

• When a student mixes 50 mL of a 1.0 M HCl and 50mL of a 1.0M NaOH in a coffee cup calorimeter, the temperature of the resultant solution increases from 21.0 ⁰C to 27.5⁰C. Calculate the enthalpy change for the reaction, assuming that the calorimeter loses only a negligible quantity of heat, that the total volume of the solution is 100 mL , that its density is 1.0 g/mL and that its specific heat is 4.18 J/g-K.

Practice Exercise

• When 50 mL of a 0.100M AgNO3 and 50mL of a 0.100M HCl are mixed in a constant –pressure calorimeter, the temperature of the mixture increases from 22.30⁰C to 23.11⁰C. The temperature increase is caused by the following reaction:

• AgNO3(aq) + HCl(aq) → AgCl(s) + HNO3(aq)Calculate ΔH for this reaction, assuming that the combined solution has a mass of 100 g and a specific heat of 4.18 J/g-K

Bomb Calorimetry: Constant –Volume Calorimetry

• 76.)One of the most important types of reactions studied using calorimetry is __combustion_____ in which a compound reacts completely with excess oxygen.

77.)Combustion reactions are most conveniently studied using a bomb calorimeter.

•78.)The substance to be studied is placed in a small cup within a sealed vessel called a _bomb__ , which is designed to withstand high pressures.

•79.)__Heat___ is released when combustion occurs.

•80.)The heat is absorbed by the calorimeter contents causing a rise in the _temperature______ of the water.

•81.)The temperature of the water is carefully measured __before____ and _after__ reaction.

•82.)To calculate the heat of combustion from the measured temperature increase in the bomb calorimeter, we must know the heat capacity of the calorimeter ,Ccal.

• Example Problem• The combustion of 1 gram of benzoic acid is known to produce 26.38 KJ of heat. If exactly 1 gram of benzoic acid is combusted in a bomb calorimeter and it increases the temperature by 4.857ᵒC, what is the heat capacity for that calorimeter?

•  • Ccal = 26.38 KJ / 4.857 ᵒC = 5.431 KJ/ᵒC

•83.)To find the heat of a reaction, qrxn using a specific calorimeter:

•Qrxn = -Ccal x ΔT

 Sample exercise 5.8

• Methylhydrazine (CH6N2) is commonly used as a liquid rocket fuel. The combustion of methylhydrazine with oxygen produces N2(g), CO2(g) and H2O(l):

• 2 CH6N2(l) + 5O2(g) → 2N2(g) + 2 CO2(g) +6 H2O(l)

• When 4.00 g of methylhydrazine is combusted in a bomb calorimeter, the temperature of the calorimeter increases from 25.00ᵒC to 39.50 ᵒC. In a separate experiment the heat capacity of the calorimeter is measured to be 7.794KJ/g-ᵒC. What is the heat of reaction for the combustion of one mole of CH6N2 in this calorimeter?

 Practice exercise

• A 0.58665g sample of lactic acid (HC3H5O3) is burned in a calorimeter whose heat capacity is 4.812 KJ/ᵒC . The temperature increases from 23.10ᵒC to 24.95ᵒC. Calculate the heat of combustion of (a) lactic acid per gram (b) per mole.

• 84.Because reactions in a bomb calorimeter are carried out under constant volume conditions, the heat transferred corresponds to a change in internal energy ΔE, rather than __enthalpy____ ΔH. However both ΔE and ΔH are usually very close in value.

Section 6 Hess’s Law

•85.)It is possible to calculate ΔH for a reaction from the __tabulated___ ΔH values of other reactions.

If a particular reaction can be carried out in one step or in a series of steps the sum of the enthalpy changes associated with the individual steps must be the same as the enthalpy change associated with the one step process. 

 CH4(g) + 2O2(g) → CO2(g) +2 H2O(g) ΔH= -802 KJ

2H2O(g) → 2H2O(l) ΔH= -88KJCH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH= -890 KJ

To obtain the net equation

1.)The sum of the reactants of the two equations is placed on the left side of the arrow.2.)The sum of all products is placed on the product side3.)Any substance that is on both sides of the equation in the same phase can be canceled like in algebraic equations.

Hess’s Law:

• If a reaction is carried out in a series of steps, ΔH for the reaction will equal the sum of the enthalpy changes for the individual steps.

• The overall enthalpy change for the process is independent of the number of steps or path in which the reaction occurs.

•  If a reaction is reversed, change the sign of Δ H.

• If a reaction is multiplied by a factor, multiply the enthalpy change by that same factor.

•85.)Hess’s Law provides a useful means of calculating _energy changes that are difficult to measure directly.

The enthalpy of combustion of C to CO2 is -393.5kJ/mol carbon, and the enthalpy of combustion of CO to CO2 is -283.0 kJ/mol CO:

C(s) + O2(g) → CO2(g) ΔH= -393.5kJCO2(g) + ½ O2(g) → CO2(g) ΔH = -283.0 kJ

Using the above data , calculate the enthalpy of combustion of C to CO:C(s) + ½ O2 → CO(g)

Practice Exercise• Carbon occurs in two forms, graphite and diamond. The

enthalpy of combustion of graphite is -393.5KJ /mol and that of diamond is -395.4kJ/mol:

• C(graphite) + O2(g) → CO2(g) ΔH= -393.5 kJ• C(diamond) + O2(g) → CO2(g) ΔH= -395.4 kJ• Calculate ΔH for the conversion of graphite to diamond:• C(graphite) → C(diamond)

 Sample Exercise 5.10• Calculate ΔH for the reaction• 2C(s) + H2(g) → C2H2(g)• given the following reactions and their respective

enthalpy changes:C2H2(g) + 5/2 O2(g) →2CO2(g) + H2O(l) ΔH=-1299.6 kJ C(s) + O2(g) → CO2(g) ΔH = -393.5 kJ H2(g) + ½ O2(g) → H2O(l)

ΔH = -285.8 kJ

Calculate ΔH for the reactionNO(g) + O(g) → NO2 (g)Given the following information:NO(g) + O3(g) → NO2(g) + O2(g) ΔH= -198.9 kJ O3(g) → 3/2 O2(g) ΔH= -142.3 kJ O2(g) → 2 O(g) ΔH= 495.0 kJ

We always get the same value of ΔH for an overall reaction regardless of how many steps we use to get to the final product.

We can use Hess’s Law to find the enthalpy change for a large number of reactions from the tabulated ∆H values.

86.)Tables exist for the following:

a.) Enthalpy of vaporization- ΔH for converting liquids to gasesb.) Enthalpy of fusion- ΔH for melting solids

86.)Tables exist for the following:

c.) Enthalpy of combustion- ΔH for combusting a substance with oxygend.) Enthalpy of formation- ΔH for formations of a substance from its elements in standard state.

∆H is dependent on temperature and pressure; so, we define the conditions at standard state.

87.)What is standard state for most enthalpies?•1 atmosphere of pressure and 25 ⁰ C

88.)What is the standard enthalpy of formation?

• ΔHf ⁰ the change in enthalpy for the

reaction that forms 1 mole of the compound from its elements with all substances in their standard state

• The stoichiometry is always arranged so that 1 mole of the product is formed.

•  

Sample exercise 5.11For which of the following reactions at 25◦C would the enthalpy change represent a standard enthalpy of formation? For those where it does not, what changes would need to be made in the reaction conditions?

a.)2Na(s) + ½ O2(g) → Na2O(s)

 b.)2K(l) + Cl2(g) → 2KCl(s)

c) C6H12O6(s) → 6C(diamond) + 6H2(g) + 3O2(g)

Practice ExerciseWrite the equation corresponding to the standard enthalpy of formation of liquid CCl4 (Carbon tetrachloride.)

We can use ∆Hf◦ of different compounds to calculate ∆H

for a reaction at standard state.

Determine ∆H◦ for the following reaction.C3H8 (g) + 5O2(g) → 3 CO2 (g) + 4 H2O(l)

On the bottom line:∆Hrxn

◦ = Σ n ∆Hf◦product - Σ m ∆Hf

◦reactants

Example 5.12Calculate the standard enthalpy of change for the combustion of 1 mole of benzene (C6H6) to CO2 (g) and H2O (l)

Practice Given the following standard enthalpy of reaction, use the standard enthalpies of formation in table 5.3 to calculate the standard enthalpy of formation of CuO (s)

Food and Fuels Section 8

•89.)What is a fuel value?•The energy released when 1 gram of a material is combusted.

The fuel value for any food or fuel can be measured by calorimetry.

•Protein has 17 kJ/g•Carbohydrates 17 kJ/g •Fats 38 kJ/g

90. Fats are the major energy storage molecules in the body for two reasons: 

•1. insoluble in water•2. produce more energy per gram than carbohydrates or proteins.

Read exercise 5.14 page 182

• A plant such as celery contains carbohydrates in the form of starch and cellulose. These two kinds of carbohydrates have essentially the same fuel values when combusted in a bomb calorimeter. When we consume celery, however, our bodies receive fuel values from the starch only. What can we conclude about the starch stored as cellulose as foods?

• A 28 –g ( 1 oz) serving of a popular breakfast cereal served with 120 mL of skim milk provides 8 g of protein, 26 g of carbohydrates and 2 g of fat. A) Using the average fuel values of these kinds of substances, estimate the amount of food energy in this serving. B.) A person of average weight uses about 100 Cal/mile when running or jogging. How many servings of this cereal provide the fuel value requirement for running 3 miles.

• Practice Exercise • Dry red beans contain 62% carbohydrate, 22% protein, and 1.5% fat. Estimate the fuel value of these beans. B.) Very light activity like reading or watching television uses about 7 kJ/min. How many minutes of such activity can be sustained by the energy provided by a serving of chicken noodle soup containing 13 g or protein, 15 g carbohydrate and 5 g of fat.

Fuels:The greater the % of carbon and hydrogen in a fuel the higher its fuel value

• Fossil fuels- coal, petroleum and natural gas• Natural gas- gaseous hydrocarbons (primarily

methane) ethane, propane and butane.• Petroleum – a liquid composed of hundreds of

compounds most of which are hydrocarbons containing sulfur, nitrogen and oxygen

• Coal- solids containing hydrocarbons of high sulfur molecular weight as well as compounds containing sulfur, oxygen and nitrogen.

Compare the quantity of heat produced by the combustion of 1 gram of propane compared to 1 gram of benzene.

Exercise 5.13The standard enthalpy change for the reaction CaCO3(s) → CaO(s) + CO2(g)

Is 178.1 KJ. From the values for the standard enthalpies of formation of CaO(s) and CO2(g) given in table 5.3 calculate the enthalpy of formation of CaCO3(s).